Chapter 5 for Chemistry

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Chapter 5
Atomic Structure and The Periodic Table
Just how small is an atom?
http://imagecache5.art.com/p/LRG/6/667/USYC000Z/fedex-field--washington-d-c-.jpg
 Has anyone been to a
professional football
stadium or a major college
football stadium?
So then, most of the atom is
just “empty space.”
 If the nucleus of an atom
was the size of a marble,
sitting at the 50 yard line, the
electrons would be about the
size of really little gnats
(bugs) whizzing around the
top rows of the upper deck.
Angstoms (Å)
2,898,550,725 U atoms in 1m!
http://intro.chem.okstate.edu/1314F00/Lecture/Chapter7/ATRADIID.DIR_PICT0003.gif

Even the largest atoms are very small. The diameter of a uranium atom is only about
0.345 nanometers.

A special unit is sometimes used to describe atomic dimensions, such as atomic radius
or atomic diameter. Note the trend as you go across a row and down a column.

That is the Angstrom. We use a Å to represent Angstroms (if you want to type that it’s
shift-alt-A on a Mac and control-shift-2, shift-A on a bogus, inferior, Windows or
Vista based machine).
Angstoms (Å)
http://upload.wikimedia.org/wikipedia/commons/1/11/Hydrogen_Atom.jpg
 Even the largest atoms are very
small. The diameter of a
uranium atom is only about
0.345 nanometers.
 0.345 nm = 3.45Å
 1nm = 10Å
 1Å = 1 x 10-10 meters
 A hydrogen atom is the smallest
atom. H has a diameter of only
0.74Å. About 13.5 billion
hydrogen atoms could fit onto
the edge of a meter stick.
What does an atom look like?
http://www.lanl.gov/orgs/pa/newsbulletin/images/Isotopes_logo.jpg
 In your notes, draw a
simple picture of an
atom. How about
Lithium.
 What did you draw?
AAA baseball club Albuquerque Isotopes logo
(you need to know what isotopes are!)
Atoms
http://upload.wikimedia.org/wikipedia/commons/thumb/e/e1/Stylised_Lithium_Atom.svg/270pxStylised_Lithium_Atom.svg.png
http://www.solarsystempictures.net/
Neutron
Proton
Electron
 Most people probably drew a
nucleus of some type with
electrons orbiting around it.
Lithium
Planets=Electrons
Sun=Nucleus
 Possibly it looks a little like a
mini solar system.
 Atoms are composed of
 Protons
 Neutrons
 Electrons
Subatomic Particles
Particle
Relative charge Relative mass Actual Mass of
(1 amu = mass of a Particle
proton)
Proton
Neutron
+1
1 amu
1.67 x 10-24g
0
1 amu
1.67 x 10-24g
-1
0 amu
9.11 x 10-28g
Electron
Electrons
 Electrons were the first particle discovered.
 JJ Thompson discovered the electron and developed the
Plum Pudding Model.
 The electron has a very small mass. It is actually
1/1840th the size of a proton (0.00054 times). In other
words, if an electron was a smidge over a pound,
protons and neutrons would weigh 1,840 POUNDS!
Protons
http://web.buddyproject.org/web017/web017/images/atom.JPG
 Protons were the next subatomic
particle to be discovered.
 Protons were discovered by E.
Goldstein.
 Protons have a relative mass of 1
and a charge of +1.
Protons
http://www.periodictable.com/Items/020.6/index.html
Atomic Number
 Protons determine the “identity” of
an atom. The number of protons is a
property called “atomic number.”
Atomic numbers are on the periodic
table.
 The number of protons determines
what kind of atom it is.
 H has 1 proton
 C has 6 protons
 U has 92 protons
Protons
 Neutral atoms have the same number of protons
and electrons. (Makes sense, right?)
 The charges balance each other out. (Ca has 20
protons, and must have 20 negatively-charged
electrons to balance out those positive charges)
 Protons are located in the nucleus of the atom.
(Where are the electrons?)
Neutrons
http://www.ct.infn.it/~rivel/Archivio/chadwick.jpg
http://kwisp.files.wordpress.com/2009/05/adventures-jimmy-neutron-300-032707.jpg
 Neutrons are also located in the nucleus
of the atom.
Ooops, wrong
neutron!
Chadwick
 The nucleus was discovered by Ernest
Rutherford, a former student of JJ
Thompson.
 The neutron was the last particle
discovered, by James Chadwick, a
former student of Rutherford.
Neutrons
 More than likely, the fact that neutrons had no charge
made it harder to discover.
 The neutron has a relative mass of 1, the same as a
proton.
 However, it has no charge. Therefore, we say that the
charge = 0.
 The actual mass of a neutron is almost the same as that of
a proton. It is slightly different
 P = 0.0000000000000000000000016726 g
 N = 0.0000000000000000000000016749 g
Golf Balls In Beakers
My Little Model of the Atom
http://www.vias.org/physics/bk4_03_02.html
 These are in the lab somewhere. Find them.
 The pink balls represent protons.
 The white balls represent neutrons.
 Scientists quickly figured out by experimentation
how many protons each element had. If you want
to read more, check out the link above. Basically
though….
The Nucleus
http://www.chemicalelements.com/bohr/b0019.gif
 Since the neutrons are located in the
nucleus, with the protons, substantially
ALL of the mass of the atom is contained
within the nucleus.
 Mass of nucleus in diagram
0.0000000000000000000000651 g
 Mass of electrons
0.0000000000000000000000000173 g
What element
is this??
 In other words, if the nucleus weighed 651
pounds, the electrons (combined) would
weigh less than a McD’s quarter-pounder
patty.
Why do the protons stick
together in the nucleus?
http://www.antonine-education.co.uk/Physics_AS/Module_1/Topic_5/strong_force.jpg
 The answer is strong nuclear force.
 It’s the strongest known force in the
universe. It far, far stronger than gravity.
 It only can be felt when the particles are
extremely close together, like when they
are packed together in the nucleus.
The secret’s in
the attractions
between the
quarks…
 Protons and neutrons are made of quarks.
It’s thought that the quarks attract other
quarks and hold the nucleus together, even
though all of the protons are positively
charged and would otherwise repel each
other.
Objective E
 We already know that the number of
protons is what makes an atom unique.
 Hydrogen has 1 proton.
 Carbon has 6 protons.
 Uranium has 92 protons.
So, if “ProtonMan”
 The “atomic number” is the number of
was a superhero,
protons. We sometimes use a Z to
he’d have a “Z” on
represent atomic number.
his suit??
Objective E
 So, for hydrogen, Z = 1
Don’t memorize
these…they are on the
Periodic Table
 For carbon, Z = 6
 For uranium, Z = 92.
 What is the atomic number for
 Aluminum
 Zinc
 Chlorine
Find THEM!!
Objective F
 So, Z (atomic number) tells us how many
protons an atom has. It does NOT tell
you how many ELECTRONS you have
(accurately) all the time!
 Most atoms have no charge. That means
that the number of protons (which are
positively charged) must balance out the
number of electrons (which are negatively
charged).
Objective F
 Unless you are TOLD that the atom has a charge, you
should assume it has no charge, and therefore, # of
protons = # of electrons.




Hydrogen (Z = 1) also has one electron.
Lithium (Z = 3) also has 3 electrons.
Carbon (Z = 6) also has 6 electrons.
Uranium (Z =92) also has 92 electrons.
 BUT REMEMBER The numbers of protons doesn’t
always equal the numbers of electrons.
Objective F
 Some atoms can lose electrons. When they do so,
they will form a positive “ion.” Some atoms can
gain electrons. When they do so, they will form a
negative “ion.”
 An ion is a atom which has an electrical charge
(either positive or negative). We’ll get to those in
Chapter 6. Na+1 and Cl-1 are formulas for ions.
 The number of protons cannot change. If the
number of protons changes, it’s no longer the same
element. Atoms can gain or lose electrons, but
they can NOT gain or lose protons in any chemical
reaction.
Schwartz’s Law
(a law I made up…hey, it’s my class)
 To calculate the number of electrons, use
 # of Electrons = Z – IC
 Where Z = atomic number and IC = ionic charge.
 Ex: Suppose we have a sodium ion with a + 1
charge. How many electrons does it have? Atomic
number (Z) is 11 (find this on Periodic Table) and
ionic charge is 1.
 # electrons = 11 - 1 = 10
Schwartz’s Law
 Let’s calculate a couple more…
 Ex: Suppose we have a sulfur ion with a - 2 charge.
How many electrons does it have? Atomic number (Z)
is 16 and ionic charge is -2.
 # electrons = 16 - (-2) = 16 + 2 = 18
 Ex: Suppose we have an zinc atom with no charge.
How many electrons does it have? Atomic number (Z)
is 30 and ionic charge is 0.
 # electrons = 30 - 0 = 30
6 neutrons
Objective F
http://www.atomicarchive.com/Physics/Images/isotopes.jpg
8 neutrons
 How do we calculate how many
neutrons we have?
 In order to do that, we need to
look at another property, called
atomic mass. The atomic mass of
an atom = THE SUM of protons
and neutrons.
 We will use another formula
Hey these are isotopes
again. Isotope = same #
of protons but a different
# of neutrons.
 # Neutrons = A – Z
 A = Mass Number
 So, what is Z again?
Objective F
http://www.lbl.gov/abc/Basic.html#Nuclearstructure
 Let’s look at an example. An atom of
Bromine (Br-80) has Z = 35 and Mass
Number = 80. How many neutrons does it
have? (Br-80 doesn’t mean bromine with a charge
of -80. When they write it like that, it’s a DASH
and 80 is the mass number)
 # Neutrons = Mass Number - Atomic
Number
 # Neutrons = 80 - 35 = 45
Objective G
http://www.usagold.com/images/gold-coins-images.jpeg
http://finestimaginary.com/shop/images/medium/jewellery/au_MED.jpg
 How do isotopes differ from each
other? (You should know this by
now).
 Look at gold (Au) on the periodic
table. It says that the mass =
196.967. Since mass number and
atomic number are ALWAYS
whole numbers, how do we get
.967?
 The answer is that the atomic
masses on the periodic table are
averages.
Objective G
 They get that average atomic mass for Au by taking
into account ALL of gold’s isotopes.
 Isotopes differ from each other in the number of
neutrons. They behave the same CHEMICALLY
because all isotopes of the same element have the
same number of protons.
 The study guide talks about water and heavy water.
Heavy water is the same as water, except that
instead of H-1 it is formed using H-2, which is
sometimes called deuterium. So while water has a
molar mass of 18, deuterium oxide or heavy water
has a mass of 20 (2 + 2 + 16)
Math Alert
Objective H
 How do we calculate the average
atomic mass?
 To do so, you need to know 2 things:
 All possible isotopes for an element
 The percent abundance for each (in
other words, how much of the whole is
represented by each isotope).
Objective H
 Let’s look at an example:
 Chlorine has 2 isotopes
 35Cl which is 75.77% of the total
amount of chlorine.
 37Cl which is 24.23% of the total
amount of chlorine.
 What is the average atomic mass of
Chlorine?
Objective H
 Cl-35 accounts for 75.77% of the total chlorine. CL-37
accounts for the rest.
 Remember to convert percents into decimals, you have to
move the decimal point 2 places to the left.
 You then mutiply the percentage (in decimal form) times the
mass number for that isotope. You do that for the other
isotope too, and then add the answers together.
 Avg Atomic Mass = 35 (0.7577) + 37 (0.2423).
 Avg Atomic Mass = 26.52 + 8.97 = 35.49
Objective H
 In our class, we are always going to round
average atomic masses to 1 decimal place.
 So, we’ll round 35.49 to 35.5 and that’s the
average atomic mass of Chlorine that we’ll
use.
 Why can’t you just average 35 and 37 (the
two isotopes) and get 36 as the average
atomic mass? Why is that wrong?
The End
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